It is always possible to increase a rate constant and accelerate a reaction by increasing the temperature. For reactions with an activation energy of 12 to 13 kcal, at temperatures around 298⁰K the rate constant doubles with every 10⁰ rise in temperature. (Can you prove this?)

But as we have seen with NH₃, there can be difficulties: The reverse reaction may be accelerated faster than the forward reaction, so that fewer products are obtained. The products or reactants may be unstable at elevated temperatures, or in special applications the surroundings may preclude the use of higher temperatures.

For example, one cannot light a match to burn glucose in the human body; this reaction must be carried out at approximately 98.6⁰F. It is for such reactions that catalysis becomes useful.

In general, catalysts lower the activation barrier for a reaction (Ea) thereby making the rate constants larger and the reactions faster. This is represented schematically by the drawing opposite.

Lowering E_a means finding an alternative pathway or mechanism for the reaction, in which the intermediate states (activated complexes) at all times are at a lower energy.

Both the forward and the reverse reactions are speeded up by a catalyst, since lowering the forward E_a necessitates lowering the reverse E_a by the same amount. A catalyst has no effect on K_eq or on the ultimate equilibrium conditions for a reaction; it only provides a way in which a spontaneous but slow reaction can arrive at equilibrium faster.

If the reaction is not already thermodynamically spontaneous, a cataIyst will be of no use. Thermodynamics does not tell a chemist how to find a catalyst for a given reaction, but it does tell him when it is, or is not, worth his time to look for one.